What are acids and bases?
In pure water at 25 °C only about two molecules in every billion have split into ions, and that tiny amount is where acids and bases begin.
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Vinegar, lemon juice and the acid in your stomach taste or act sour and sharp. Soap, baking soda and oven cleaner feel slippery and bitter. People sorted substances into these two families long before anyone knew about atoms. In 1680 Robert Boyle described acids as solutions that dissolve many substances, change the colours of some natural dyes and lose these powers when they meet an alkali. Chemists in the eighteenth century added that acids react with limestone to give off a gas, now known to be carbon dioxide, and that acids and alkalis together make neutral substances.
From hydrogen to hydronium
In 1815 Humphry Davy showed that hydrogen, not oxygen, is the essential ingredient of acids. In 1884 Svante Arrhenius gave the first modern definition: an acid is a compound that dissolves in water to give hydrogen ions, and a base is a compound that dissolves in water to give hydroxide ions, OH⁻. A bare hydrogen ion is just a proton, and in water it never stays bare. It attaches to a water molecule to form the hydronium ion, H₃O⁺. When hydrogen chloride gas dissolves, the solution, hydrochloric acid, is full of H₃O⁺ and Cl⁻ ions.
Acids give protons, bases take them
In 1923 Johannes Brønsted and Thomas Lowry proposed a broader idea that chemists still use. An acid is a proton donor and a base is a proton acceptor, and an acid-base reaction is simply a proton moving from one to the other. When hydrogen fluoride dissolves, HF gives a proton to H₂O, making H₃O⁺ and F⁻. When ammonia dissolves, NH₃ takes a proton from H₂O, making NH₄⁺ and OH⁻. Water acts as a base in the first reaction and as an acid in the second. Substances that can do both are called amphiprotic.
After an acid gives up its proton, what is left can take a proton back, so it is a base. Chemists call it the conjugate base of the acid. F⁻ is the conjugate base of HF, and Cl⁻ is the conjugate base of HCl. Going the other way, NH₄⁺ is the conjugate acid of ammonia. Every Brønsted-Lowry reaction has two such pairs, each linking an acid to its conjugate base across the reaction arrow.
Water talks to itself
Because water is amphiprotic, two water molecules can trade a proton with each other: H₂O + H₂O ⇌ H₃O⁺ + OH⁻. This autoionization happens only slightly. At 25 °C about two out of every billion water molecules are ionized, giving 1.0 × 10⁻⁷ mol of H₃O⁺ and the same amount of OH⁻ in every litre. Their product, the ion-product constant of water, is Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴. Add an acid and [H₃O⁺] goes up while [OH⁻] goes down, always keeping the product at Kw. That one equation underlies the whole pH scale in the next unit.
- 1680: Boyle lists the properties of acids
- 1815: Davy: hydrogen is essential to acids
- 1884: Arrhenius: acids give H⁺, bases give OH⁻ in water
- 1923: Brønsted and Lowry: proton donors and acceptors
- In the Brønsted-Lowry definition, an acid is…
- A proton donor
- A proton acceptor
- An electron donor
- Any salt
- What is the formula of the hydronium ion?
- H₃O⁺
- OH⁻
- H₂O₂
- H⁻
- What is the conjugate base of HF?
- F⁻
- H₃O⁺
- HF₂
- OH⁻
- When NH₃ dissolves in water, which ion makes the solution basic?
- OH⁻
- H₃O⁺
- Cl⁻
- Na⁺
- A substance that can act as both an acid and a base is called…
- Amphiprotic
- Neutral
- Conjugate
- Insoluble
- In which year did Arrhenius define acids and bases?
- What is Kw at 25 °C?
- 1.0 × 10⁻¹⁴
- 1.0 × 10⁻⁷
- 7
- 14
- Balance Mg + HCl → MgCl₂ + H₂. What is the coefficient of HCl?
- Which gas forms when an acid reacts with limestone (CaCO₃)?
- Carbon dioxide
- Hydrogen
- Oxygen
- Chlorine
- In pure water at 25 °C, what is [OH⁻] in mol/L?
- 1.0 × 10⁻⁷
- 1.0 × 10⁻¹⁴
- 0
- 1.0
- Arrhenius (1884): acids give H⁺, bases give OH⁻ in water. Brønsted-Lowry (1923): acids donate protons, bases accept them.
- In water H⁺ becomes H₃O⁺; an acid minus H⁺ is its conjugate base.
- Water autoionizes slightly: Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C.
Sources for this course (10)
- OpenStax Chemistry 2e (openstax.org/books/chemistry-2e): 4.2 neutralization and antacids (stomach roughly 0.03 M HCl, milk of magnesia Mg(OH)2, Al(OH)3)
- 14.1 Brønsted-Lowry acids and bases (Boyle 1680, Davy 1815, Arrhenius 1884, Brønsted and Lowry 1923, autoionization about two in a billion molecules, Kw = 1.0 × 10⁻¹⁴ at 25 °C, about 5.6 × 10⁻¹³ at 100 °C)
- 14.2 pH and pOH (pH = −log[H3O+], stomach acid 1.2 × 10⁻³ M gives 2.92, vinegar about 4 × 10⁻³ M gives 2.4, blood 7.3, KOH 0.0125 M gives 12.10, neutral pH 6.31 at 80 °C, rainwater 5 to 6, acid rain below 5, pH meter resolution 0.001, Figure 14.2: 1 M HCl 0, gastric juice about 1.5, lime juice about 2, wine about 3.5, coffee about 5, milk of magnesia about 10.5)
- 14.3 relative strengths (six strong acids and six strong bases, Ka, percent ionization, 0.100 M acetic acid 1.3%)
- 14.6 buffers (normal blood pH about 7.4)
- 14.7 titrations (25.00 mL of 0.100 M HCl or acetic acid with 0.100 M NaOH, Table 14.2, equivalence pH 7.00 and 8.72, half-equivalence pH = pKa, Ka acetic acid 1.8 × 10⁻⁵, methyl orange Ka 4.0 × 10⁻⁴ red HIn and yellow In⁻, interval about pKa ± 1, phenolphthalein colourless below pH 8.3)
- LibreTexts Acid-Base Indicators (litmus about 5 to 8, methyl orange 3.1 to 4.4, phenolphthalein 8.3 to 10.0; chem.libretexts.org) and bromothymol blue 6.0 to 7.6 yellow to blue (Prairie South indicator table via LibreTexts Chemistry 30)
- PubChem acetic acid pKa 4.76 at 25 °C, sodium bicarbonate 0.1 M pH 8.3, ammonia 0.1 N pH 11.1, sodium hydroxide (pubchem.ncbi.nlm.nih.gov)
- NOAA Ocean acidification (noaa.gov: surface ocean pH has fallen by 0.1 since the industrial revolution, about a 30 percent increase in acidity)
- IUPAC CIAAW abridged standard atomic weights.
Checked: 2026-10-17